Question

The Group 5 A elements can form molecules or ions that involve three, five, or six covalent bonds; \(\mathrm{NH}_{3}, \mathrm{AsCl}_{5}\), and \(\mathrm{PF}_{6}^{-}\) are examples. Draw the Lewis structure for each of these substances, and predict the molecular structure and hybridization for each. Why doesn't \(\mathrm{NF}_{5}\) or \(\mathrm{NCl}_{6}^{-}\) form?

Short answer

The Lewis structures, molecular structures, and hybridizations for NH₃, AsCl₅, and PF₆⁻ are as follows: 1. NH₃: Lewis structure: N with 3 single bonds to H atoms and 1 lone pair. Molecular structure: Trigonal pyramidal. Hybridization: sp³. 2. AsCl₅: Lewis structure: As with 5 single bonds to Cl atoms. Molecular structure: Trigonal bipyramidal. Hybridization: sp³d. 3. PF₆⁻: Lewis structure: P with 6 single bonds to F atoms. Molecular structure: Octahedral. Hybridization: sp³d². NF₅ and NCl₆⁻ do not form because nitrogen's 2p orbitals cannot accommodate enough electrons for the required number of bonds and due to steric hindrance from the larger Cl atoms.

Step-by-step solution

NH₃ (Ammonia)

1. Calculate the total number of valence electrons: N has 5 valence electrons and H has 1 valence electron. Since there are 3 H atoms, the total number of valence electrons is 5 + (3 × 1) = 8. 2. Draw the Lewis structure: Place N in the center, and bond each H to N with a single bond. N will have 1 lone pair of electrons remaining. H | H - N | H 3. Determine the molecular structure and hybridization: As there are 3 bonded atom groups and 1 lone pair, the electron pair geometry is tetrahedral. Due to the presence of a lone pair, the molecular structure is trigonal pyramidal. Hybridization is sp³.

AsCl₅ (Arsenic pentafluoride)

1. Calculate the total number of valence electrons: As has 5 valence electrons and Cl has 7 valence electrons. Since there are 5 Cl atoms, the total number of valence electrons is 5 + (5 × 7) = 40. 2. Draw the Lewis structure: Place As in the center, and bond each Cl to As with a single bond. Cl - As - Cl | _ _ | Cl - - Cl | Cl 3. Determine the molecular structure and hybridization: As there are 5 bonded atom groups and no lone pairs, the electron pair geometry and molecular structure are trigonal bipyramidal. Hybridization is sp³d.

PF₆⁻ (Hexafluorophosphate)

1. Calculate the total number of valence electrons: P has 5 valence electrons and F has 7 valence electrons. Since there are 6 F atoms, the total number of valence electrons is 5 + (6 × 7) + 1 (due to the negative charge) = 48. 2. Draw the Lewis structure: Place P in the center, and bond each F to P with a single bond. F | F - P - F _|_| F F | F 3. Determine the molecular structure and hybridization: As there are 6 bonded atom groups and no lone pairs, the electron pair geometry and molecular structure are octahedral. Hybridization is sp³d².

Explaining why NF₅ and NCl₆⁻ do not form

NF₅ cannot form because nitrogen has only 5 valence electrons. For 5 single bonds with F, it would need 10 electrons, which would require borrowing more electrons than it can accommodate in its 2p orbitals. NCl₆⁻ cannot form because nitrogen has only 5 valence electrons, and for 6 single bonds with Cl, it would need 12 electrons. This would also require borrowing more electrons than it can accommodate in its 2p orbitals. Additionally, the bulky size of the Cl atoms may create repulsive forces, preventing them from forming such a structure.

Key concepts

Molecular Geometry

Molecular geometry refers to the three-dimensional arrangement of atoms within a molecule. Understanding this geometry aids in predicting the physical and chemical properties of the molecule. Each molecule can adopt different shapes depending on the number of bonds and lone electron pairs around the central atom.

For example, in ammonia (\( ext{NH}_3\)), there are three hydrogen atoms bonded to a central nitrogen atom and one lone pair of electrons. This leads to a molecular geometry known as trigonal pyramidal. Such a shape is slightly distorted from a perfect tetrahedral geometry due to lone pair-bond pair repulsion, which pushes the hydrogen atoms slightly closer together.

In contrast, arsenic pentafluoride (\( ext{AsCl}_5\)) adopts a different shape because it has five chlorine atoms and no lone pairs on the arsenic. This results in a trigonal bipyramidal geometry, where the atoms arrange themselves in two triangle-like planes. Similarly, hexafluorophosphate (\( ext{PF}_6^-\)) has six fluorine atoms around phosphorus with an octahedral geometry, symmetric around the central atom. Understanding these shapes is essential for grasping the spatial distribution of molecules and predicting how they interact with others.

Electron Pair Geometry

Electron pair geometry is the spatial arrangement of all electron pairs around the central atom, both bonding and non-bonding. This concept helps predict molecular geometry but considers all electron clouds equally, rather than just focusing on bonded atoms.

Take ammonia (\( ext{NH}_3\)), for instance. Although its molecular shape is trigonal pyramidal, its electron pair geometry is tetrahedral. This is due to nitrogen being surrounded by four areas of electron density (three bonds and one lone pair), leading to a tetrahedral arrangement when considering the electron pairs.

In arsenic pentafluoride (\( ext{AsCl}_5\)), the electron pair geometry is trigonal bipyramidal, matching its molecular shape as there are no lone pairs. For hexafluorophosphate (\( ext{PF}_6^-\)), with its six fluorine atoms, the electron pair geometry is octahedral, reflecting the six regions of electron density. By understanding electron pair geometry, one can see how both bonding and non-bonding electrons dictate the overall molecular shape.

Hybridization

Hybridization is a concept used to explain the mixing of atomic orbitals to form new hybrid orbitals that are suitable for the pairing of electrons to form chemical bonds. It provides a simple model that explains molecular geometries derived from valence bond theory.

In ammonia (\( ext{NH}_3\)), the nitrogen atom undergoes sp³ hybridization. That means one s and three p orbitals mix to form four equivalent sp³ hybrid orbitals, leading to the tetrahedral arrangement associated with this type of hybridization.

Arsenic in arsenic pentafluoride (\( ext{AsCl}_5\)) undergoes sp³d hybridization. This involves one s, three p, and one d orbital combining to accommodate the five bonds, aligning with the trigonal bipyramidal molecular geometry. Similarly, phosphorus in hexafluorophosphate (\( ext{PF}_6^-\)) is sp³d² hybridized. This includes six hybrid orbitals formed from the mixing of one s, three p, and two d orbitals, resulting in an octahedral geometry. Understanding hybridization is key to predicting and explaining molecular shapes and bond angles.

Valence Electrons

Valence electrons are the electrons in the outermost shell of an atom. They are crucial as they determine an atom's ability to bond with other atoms and influence both the formation and strength of chemical bonds.

In ammonia (\( ext{NH}_3\)), each hydrogen atom contributes one valence electron, while nitrogen has five. This totals eight valence electrons available for bonding, which constructs the molecule's structure and explains its bonding behavior.

Arsenic in arsenic pentafluoride (\( ext{AsCl}_5\)) and phosphorus in hexafluorophosphate (\( ext{PF}_6^-\)) each have five valence electrons. Coupling this with the valence electrons of the chlorine or fluorine atoms used in bonding, more complex structures are formed. In fact, understanding the limits of valence electrons explains why molecules like \( ext{NF}_5\) or \( ext{NCl}_6^−\) cannot form—nitrogen's valence shell cannot accommodate such high coordination due to its electron configuration. Valence electrons are thus fundamental elements in predicting how molecules are built and understanding why certain compounds do not exist.